11th Grade Chemistry Notes and Review

 These are notes and a review of 11th grade or high school chemistry. 11th grade chemistry covers all the material listed here, but this is a concise review of what you need to know to pass a cumulative final exam. There are several ways to organize the concepts. Here is the categorization I've chosen for these notes:

Chemical and Physical Properties and Changes

11th grade chemistry covers key topics.
11th grade chemistry covers key topics. Chris Ryan / Getty Images

Chemical Properties: properties that describe how one substance reacts with another substance. Chemical properties may only be observed by reacting one chemical with another.

Examples of Chemical Properties:

  • flammability
  • oxidation states
  • reactivity

Physical Properties: properties used to identify and characterize a substance. Physical properties tend to be ones you can observe using your senses or measure with a machine.

Examples of Physical Properties:

  • density
  • color
  • melting point

Chemical vs Physical Changes

Chemical Changes result from a chemical reaction and make a new substance.

Examples of Chemical Changes:

  • burning wood (combustion)
  • rusting of iron (oxidation)
  • cooking an egg

Physical Changes involve a change of phase or state and do not produce any new substance.

Examples of Physical Changes:

  • melting an ice cube
  • crumpling a sheet of paper
  • boiling water

Atomic and Molecular Structure

This is a diagram of a helium atom, which has 2 protons, 2 neutrons, and 2 electrons.
This is a diagram of a helium atom, which has 2 protons, 2 neutrons, and 2 electrons. Svdmolen/Jeanot, Public Domain

The building blocks of matter are atoms, which join together to form molecules or compounds. It's important to know the parts of an atom, what ions and isotopes are, and how atoms join together.

Parts of an Atom

Atoms are made up of three components:

  • protons - positive electric charge
  • neutrons - no electric charge
  • electrons - negative electric charge

Protons and neutrons form the nucleus or center of each atom. Electrons orbit the nucleus. So, the nucleus of each atom has a net positive charge, while the outer portion of the atom has a net negative charge. In chemical reactions, atoms lose, gain, or share electrons. The nucleus does not participate in ordinary chemical reactions, although nuclear decay and nuclear reactions can cause changes in the atomic nucleus.

Atoms, Ions, and Isotopes

The number of protons in an atom determines which element it is. Each element has a one- or two-letter symbol that is used to identify it in chemical formulas and reactions. The symbol for helium is He. An atom with two protons is a helium atom regardless of how many neutrons or electrons it has. An atom may have the same number of protons, neutrons, and electrons or the number of neutrons and/or electron may differ from the number of protons.

Atoms that carry a net positive or negative electric charge are ions. For example, if a helium atom loses two electrons, it would have a net charge of +2, which would be written He2+.

Varying the number of neutrons in an atom determines which isotope of an element it is. Atoms may be written with nuclear symbols to identify their isotope, where the number of nucleons (protons plus neutrons) is listed above and to the left of an element symbol, with the number of protons listed below and to the left of the symbol. For example, three isotopes of hydrogen are:

11H, 21H, 31H

Since you know the number of protons never changes for an atom of an element, isotopes more commonly are written using the element symbol and the number of nucleons. For example, you could write H-1, H-2, and H-3 for the three isotopes of hydrogen or U-236 and U-238 for two common isotopes of uranium.

Atomic Number and Atomic Weight

The atomic number of an atom identifies its element and its number of protons. The atomic weight is the number of protons plus the number of neutrons in an element (because the mass of electrons is so small compared with that of protons and neutrons that it essentially doesn't count). The atomic weight sometimes is called atomic mass or the atomic mass number. The atomic number of helium is 2. The atomic weight of helium is 4. Note that the atomic mass of an element on the periodic table isn't a whole number. For example, the atomic mass of helium is given as 4.003 rather than 4. This is because the periodic table reflects the natural abundance of isotopes of an element. In chemistry calculations, you use the atomic mass given on the periodic table, assuming a sample of an element reflects the natural range of isotopes for that element.


Atoms interact with each other, often forming chemical bonds with each other. When two or more atoms bond to each other, they form a molecule. A molecule can be simple, such as H2, or more complex, such as C6H12O6. The subscripts indicate the number of each type of atom in a molecule. The first example describes a molecule formed by two atoms of hydrogen. The second example describes a molecule formed by 6 atoms of carbon, 12 atoms of hydrogen, and 6 atoms of oxygen. While you could write the atoms in any order, the convention is to write the positively charged past of a molecule first, followed by the negatively charged part of the molecule. So, sodium chloride is written NaCl and not ClNa.

Periodic Table Notes and Review

This is the periodic table of the elements.
This is the periodic table of the elements, with different colors identifying element groups. Todd Helmenstine

The periodic table is an important tool in chemistry. These notes review the periodic table, how it is organized, and periodic table trends.

Invention and Organization of the Periodic Table

In 1869, Dmitri Mendeleev organized the chemical elements into a periodic table much like the one we use today, except his elements were ordered according to increasing atomic weight, while the modern table is organized by increasing atomic number. The way the elements are organized makes it possible to see trends in element properties and to predict the behavior of elements in chemical reactions.

Rows (moving left to right) are called periods. Elements in a period share the same highest energy level for an unexcited electron. There are more sub levels per energy level as atom size increases, so there are more elements in periods further down the table.

Columns (moving top to bottom) form the basis for element groups. Elements in groups share the same number of valence electrons or outer electron shell arrangement, which gives elements in a group several common properties. Examples of element groups are alkali metals and noble gases.

Periodic Table Trends or Periodicity

The organization of the periodic table makes it possible to see trends in properties of elements at a glance. The important trends relate to an atomic radius, ionization energy, electronegativity, and electron affinity.

  • Atomic Radius
    Atomic radius reflects the size of an atom. Atomic radius decreases moving from left to right across a period and increases moving from top to bottom down an element group. Although you might think atoms would simply become larger as they gain more electrons, electrons remain in a shell, while the increasing number of protons pulls the shells in closer to the nucleus. Moving down a group, electrons are found further from the nucleus in new energy shells, so the overall size of the atom increases.
  • Ionization Energy
    Ionization energy is the amount of energy needed to remove an electron from an ion or atom in the gas state. Ionization energy increases moving from left to right across a period and decreases moving top to bottom down a group.
  • Electronegativity
    Electronegativity is a measure of how easily an atom forms a chemical bond. The higher the electronegativity, the higher the attraction for bonding an electron. Electronegativity decreases moving down an element group. Elements on the lefthand side of the periodic table tend to be electropositive or more likely to donate an electron than accept one.
  • Electron Affinity
    Electron affinity reflects how readily an atom will accept an electron. Electron affinity varies according to element group. The noble gases have electron affinities near zero because they have filled electron shells. The halogens have high electron affinities because the addition of an electron gives an atom a completely filled electron shell.


Chemical Bonds and Bonding

This is a photograph of an ionic bond between two atoms.
This is a photograph of an ionic bond between two atoms. Wikipedia GNU Free Documentation License

Chemical bonds are easy to understand if you keep in mind the following properties of atoms and electrons:

  • Atoms seek the most stable configuration.
  • The Octet Rule states that atoms with 8 electrons in their outer orbital will be most stable.
  • Atoms can share, give, or take electrons of other atoms. These are forms of chemical bonds.
  • Bonds occur between the valence electrons of atoms, not the inner electrons.

Types of Chemical Bonds

The two main types of chemical bonds are ionic and covalent bonds, but you should be aware of several forms of bonding:

  • Ionic Bonds
    Ionic bonds form when one atom takes an electron from another atom.

    Example: NaCl is formed by an ionic bond where sodium donates its valence electron to chlorine. Chlorine is a halogen. All halogens have 7 valence electrons and need one more to gain a stable octet. Sodium is an alkali metal. All alkali metals have 1 valence electron, which they readily donate to form a bond.

  • Covalent Bonds
    Covalent bonds form when atoms share electrons. Really, the main difference is the electrons in ionic bonds are more closely associated with one atomic nucleus or the other, which electrons in a covalent bond are about equally likely to orbit one nucleus as the other. If the electron is more closely associated with one atom than the other, a polar covalent bond may form.

    Example: Covalent bonds form between hydrogen and oxygen in water, H2O.

  • Metallic Bond
    When the two atoms both are metals, a metallic bond forms. The difference in a metal is that the electrons could be any metal atom, not just two atoms in a compound.

    Example: Metallic bonds are seen in samples of pure elemental metals, such as gold or aluminum, or alloys, such as brass or bronze.

Ionic or Covalent?

You may be wondering how you can tell whether a bond is ionic or covalent. You can look at the placement of elements on the periodic table or a table of element electronegativities to predict the type of bond that will form. If the electronegativity values are very different from each other, an ionic bond will form. Usually, the cation is a metal and the anion is a nonmetal. If the elements both are metals, expect a metallic bond to form. If the electronegativity values are similar, expect a covalent bond to form. Bonds between two nonmetals are covalent bonds. Polar covalent bonds form between elements that have intermediate differences between the electronegativity values. 

How To Name Compounds - Chemistry Nomenclature

In order for chemists and other scientists to communicate with each other, a system of nomenclature or naming was agreed upon by the International Union of Pure and Applied Chemistry or IUPAC. You'll hear chemicals called their common names (e.g., salt, sugar, and baking soda), but in the lab you would use systematic names (e.g., sodium chloride, sucrose, and sodium bicarbonate). Here's a review of some key points about nomenclature.

Naming Binary Compounds

Compounds may be made up of only two elements (binary compounds) or more than two elements. Certain rules apply when naming binary compounds:

  • If one of the elements is a metal, it is named first.
  • Some metals can form more than one positive ion. It is common to state the charge on the ion using Roman numerals. For example, FeCl2 is iron(II) chloride.
  • If the second element is a nonmetal, the name of the compound is the metal name followed by a stem (abbreviation) of the nonmetal name followed by "ide". For example, NaCl is named sodium chloride.
  • For compounds consisting of two nonmetals, the more electropositive element is named first. The stem of the second element is named, followed by "ide". An example is HCl, which is hydrogen chloride.

Naming Ionic Compounds

In addition to the rules for naming binary compounds, there are additional naming conventions for ionic compounds:

  • Some polyatomic anions contain oxygen. If an element forms two oxyanions, the one with less oxygen ends in -ite while the one with more oxgyen ends in -ate. For example:
    NO2- is nitrite
    NO3- is nitrate